Chapter 13: Phenomenapeople.chem.ucsb.edu/feldwinn/darby/Chem1B/Overheads/Chapter-1… · Chapter...

Chapter 13: Bonding: General Concepts

Chapter 13: Phenomena

Phenomena: Scientists measured the bond angles of some common molecules.

In the pictures below each line represents a bond that contains 2 electrons. If

multiple lines are drawn together these are double or triple bonds and contain 4

and 6 electrons respectively. What patterns do you notice from the data?

H HN

H

H HN

H

H +

O

CO

O

CN H

Cl ClC

Cl

Cl

B

F

Bond Angles: 107˚

Bond Angles: 109.5˚ Bond Angles: 109.5˚

Bond Angles: 120˚

C

O

Bond Angles: 105˚

Bond Angles: 119˚

Bond Angles: 120˚ Bond Angle: 117˚

Bond Angles: 180˚

Bond Angles: 180˚

a) b) c)

d)

g)

e)

h)

f)

i)

j)

S

O

Chapter 13

Bonding:

General Concepts

o Types of Bonding

o Electronegativity

o Lewis Structures

o Strength/Length of

Covalent Bonds

o Shapes of Molecules (VSEPR)

o Polar Molecules

2

Big Idea: Bonds are formed from

the attraction

between oppositely

charged ions or by

sharing electrons.

Only the valence

electrons participate

in bonding. The shape

of the molecules

maximize the distance

between areas of high

electron density.


Types of Bonding

Ionic Bonds: Formed when a lower

energy can be achieved by the

complete transfer of one or more

electrons from the atoms of one element

to those of another; the compound is

then held together by electrostatic

attraction between the ions.

Covalent Bonds: Formed when the lowest

energy structure can be achieved by

sharing electrons.

3


Types of Bonding

Ionic Bonds: Tend to be between a

metal and a non metal.

Metals: Usually lose their electrons.

Nonmetals: Usually accept additional

electrons.

4

Note: In general, atoms gain or lose electrons until they have the same number of

electrons as the nearest noble gas


Types of Bonding

Element

Electron

Configuration

(Atom)

Gain/Lose

Electrons

Ion

Formed

Electron

Configuration

(Ion)

S

K

I

5

What ions do atoms form?


Types of Bonding

Ionic Solids: Assembly of cations and anions

stacked together in a regular array.

Ionic Compounds are represented with formula

units (lowest ratio of types of atoms in the

compound).

Ionic Compound

Covalent Compound

6

AB2 2AB2 3AB2

CD 2CD 3CD+ -+ -

+-

+ -+-

+ -

Note: Ionic solids are example of crystalline arrays in which the overall charge on

an ionic solid is neutral.


Types of Bonding

Steps to calculate the energy needed to form an ionic bond:

Step 1: Standard states to gaseous single atom state

Na(s) Na(g) 97 𝑘𝐽𝑚𝑜𝑙

½F2(g) F(g) 80. 𝑘𝐽

𝑚𝑜𝑙

Step 2: Both atoms have to form ions

Na(g) Na+(g) + e-(g) 494 𝑘𝐽𝑚𝑜𝑙

(ionization energy)

F(g) + e-(g) F-(g) -323 𝑘𝐽𝑚𝑜𝑙

((-)electron affinity)

Step 3: The ions need to come together to form a

crystal (Lattice Energy)

Na+(g) + F-(g) NaF(s) -923 𝑘𝐽𝑚𝑜𝑙

Total Reaction

Na(s) + ½F2(g) NaF(s)

97 𝑘𝐽

𝑚𝑜𝑙+ 90. 𝑘𝐽

𝑚𝑜𝑙+ 494 𝑘𝐽

𝑚𝑜𝑙+ −323 𝑘𝐽

𝑚𝑜𝑙+−923 𝑘𝐽

𝑚𝑜𝑙= −575 𝑘𝐽

𝑚𝑜𝑙

7

Note: When energy is

released, the sign is negative

because no work is needed to

make the reaction happen.


Types of Bonding

What is holding ionic solids together?

Coulombic Potential Energy

𝐸𝑃,12 =𝑍1𝑒 𝑍2𝑒

4𝜋𝜀∘𝑟12= +𝑍 −𝑍 𝑒2

4𝜋𝜀∘𝑑=−𝑍2𝑒2

4𝜋𝜀∘𝑑

Z1 & Z2 = charge of ions

The total potential energy is the sum of all the

potential energies

𝐸𝑃 =1

4𝜋𝜀∘−𝑍2𝑒2

𝑑+𝑍2𝑒2

2𝑑−𝑍2𝑒2

3𝑑+𝑍2𝑒2

4𝑑⋯ = − 𝑍2𝑒2

4𝜋𝜀∘𝑑1−

1

2+1

3−1

4+⋯

𝐸𝑃 = −𝑙𝑛 2𝑍2𝑒2

4𝜋𝜀∘𝑑

Need to multiply by 2 to account for the other half of the

line.

𝐸𝑃 = −2𝑙𝑛 2𝑍2𝑒2

4𝜋𝜀∘𝑑

8

Note: 𝑙𝑛 2 = 1 − 1

2+ 1

3− 1

4+⋯


Types of Bonding

Once neighboring ions come into contact they

start to repel each other.

𝐸𝑃∗ ∝ 𝑒− 𝑑 𝑑∗

d=0 𝑒− 𝑑 𝑑∗ = 1 max repulsion

d>1 𝑒− 𝑑 𝑑∗ = 1 repulsion decreases

The potential energy of an

ionic solid is a combination of

the favorable Coulombic

interaction of the ions and the

unfavorable exponential

increase which results when

the atoms touch. The ideal

bond length occurs at the

minimum potential energy.

Energy Minimum Occurs: 𝐸𝑝,𝑚𝑖𝑛 = −𝑁𝐴 𝑍1𝑍24𝜋𝜀∘𝑑

1 − 𝑑

𝑑∗𝐴

9

Note: d* is a constant

that is commonly

taken to be 34.5 pm


Types of Bonding

Covalent Bond: A pair of electrons shared

between two atoms (occurs between

two non metals)

10

Note: In covalent bond formation, atoms go as far as possible toward completing

their octets by sharing electron pairs.


Electronegativity

Electronegativity (χ): The ability of an

atom to attract electrons to itself when it

is part of a compound

11

Note: The atom with

higher electronegativity

has a stronger

attractive power on

electrons and pulls the

electrons away from the

atom with the lower

electronegativity.


Electronegativity

The dividing line between ionic and covalent bonds is hazy.

12

Difference in Electronegativity Type of Bond

> 1.8 Mostly Ionic

0.4-1.8 Polar Covalent

< 0.4 Mostly Covalent

0 Non-polar Covalent

NaCl MgCl2 AlCl3 SiCl4 PCl3 S2Cl2 Cl2ionic polar-covalent covalent


Electronegativity

What makes covalent bonds partly ionic?

Electric Dipole: A positive charge next to an

equal but opposite negative charge.

Electric Dipole Moment (𝝁): The magnitude of

the electric dipole [units debye (𝐷)].

Polar Covalent Bond: A covalent bond between

atoms that have partial electric charges.

13

Note: The dipole moment associated with H-Cl is about 1.1 𝐷.


Electronegativity

What makes ionic bonds partly covalent?

Polarizability (𝜶): The ease with which the

electron cloud of a molecule can be distorted.

As the cation’s positive

charge pulls on the

anion’s negative

electrons, the spherical

electron cloud of the

anion becomes distorted

in the direction of the

cation. This causes the

bond to have covalent

bond properties.

14

Note: The larger the anion the easier it is to distort the electron cloud.


Lewis Structures

Lewis Symbols: The chemical symbol of an

element, with a dot for each valence electron.

Step 1: Determine the number of valence e- from

electron configuration.

Step 2: Place 1 dot around the element for each

valence electron.

15


Lewis Structures

Drawing Ionic Lewis Structures

Step 1: Determine the electron configuration of

the elements in the compound.

Step 2: Determine the electron configuration of

the ions that the elements form.

Step 3: Draw the Lewis symbols for the ions. Do

not forget to include charge. The cations should

have no electrons around them and the anions

should have 8 electrons around them.

Step 4: Organize the Lewis symbols such that

cations are next to anions.

16


Lewis Structures

General Rules (Covalent Lewis Structures)

All valence electrons of the atoms in the Lewis

structures must be shown.

Generally, electrons are paired. Except for

odd electron molecules such as NO and NO2.

Generally, each atom has 8 electrons in its

valence shell with the exception of H which

only needs 2 valance electrons.

Multiple bonds (double and triple bonds) can

be formed.

Show atoms by their chemical symbols (ex. H)

Show covalent bonds by lines (ex. F–F)

Show lone pairs of electrons by pairs of dots (ex. :)

17


Lewis Structures

Drawing Covalent Lewis Structures (for structures

that obey octet rule)

Step 1: Count the number of valence electrons

on each atom; for ions adjust the number of

electrons to account for the charge.

Step 2: Calculate the number of electrons that

are needed to fill each atom’s octet (or duplet,

in the case of H).

Step 3: Calculate the number of bonds:

# 𝐵𝑜𝑛𝑑𝑠 =𝑊𝑎𝑛𝑡𝑒𝑑 𝑒− 𝑆𝑡𝑒𝑝 2 −𝑉𝑎𝑙𝑒𝑛𝑐𝑒 𝑒− 𝑆𝑡𝑒𝑝 1

2.

Step 4: Calculate the number of electrons left

over: # 𝑒− = 𝑉𝑎𝑙𝑒𝑛𝑐𝑒 𝑒− − 2 # 𝐵𝑜𝑛𝑑𝑠 .

Step 5: Place bonds/electrons around elements

so that octets/duplet are satisfied.18


Lewis Structures

Tips When Drawing Lewis Structures

How to pick central atom:

Choose the central atom to be the atom with the lowest

ionization energy (atom closest to the lowest left hand

corner of the periodic table)

Arrange the atoms symmetrically around the central

atom

19

Examples:

SO2 would be arranged OSO not SOO

Note: Acids are an exception to the rule because H is written first in acids.

Note: In simple formulas the central atom is often written first, followed by the

atoms attached to it.


Lewis Structures

What is the Lewis Structure for SCN-?

# 𝑏𝑜𝑛𝑑𝑠 =𝑤𝑎𝑛𝑡𝑒𝑑 𝑒− − 𝑣𝑎𝑙𝑒𝑛𝑐𝑒 𝑒−

2=24−16

2= 4

# 𝑒− = 𝑣𝑎𝑙𝑒𝑛𝑐𝑒 𝑒− − 2 # 𝑏𝑜𝑛𝑑𝑠 = 16 − 2 4 = 8

Which structure is most likely?

20

1(S) 1(C) 1(N) 1(e-) Total

Valence e- 1(6) 1(4) 1(5) 1 16

Wanted e- 1(8) 1(8) 1(8) 24

___


Lewis Structures

Formal Charge: The electric charge of an atom

in a molecule assigned on the assumption that

the bonding is nonpolar covalent.

Formal Charge = Valence e- – e- Surrounding Atom

Generally, compounds with the lowest formal

charges possible (charges closest to 0) are

favored.

21

Note: The formal charge on neutral molecules must add up to zero.

Note: The formal charge on ions must add up to the charge on the ion.


Lewis Structures

Formal charge and oxidation numbers both give

us information about the number of electrons

around an atom in a compound.

Formal Charge Exaggerates Covalent Character

Assumes that the electrons are shared equally by all atoms.

Oxidation Number Exaggerates Ionic Character

Assumes that octets are complete filled and all

electrons must only belong to one atom.

22

0 0 0

-2 4 -2

O●●●●

●●

●● C

2-4+

O●●●●

●●

●●

2-


Lewis Structures

Resonance: A blend of Lewis structures

into a single composite hybrid structure.

Resonance Hybrid: The composite

structure that results from a resonance.

Delocalized Electrons: Electrons that are

spread over several atoms in a molecule.

23


Lewis Structures

What is the Lewis Structure for PO43-?

# 𝑏𝑜𝑛𝑑𝑠 =𝑤𝑎𝑛𝑡𝑒𝑑 𝑒− − 𝑣𝑎𝑙𝑒𝑛𝑐𝑒 𝑒−

2=40−32

2= 4

# 𝑒− = 𝑣𝑎𝑙𝑒𝑛𝑐𝑒 𝑒− − 2 # 𝑏𝑜𝑛𝑑𝑠 = 40 − 2 4 = 32

24

1(P) 4(O) 1(e-) Total

Valence e- 1(5) 4(6) 1(3) 32

Wanted e- 1(8) 4(8) 40


Lewis Structures

When the central atom in a molecule has empty

d-orbitals, it may be able to accommodate 10,

12, or even more electrons, this is referred to as

an expanded valence shell.

Size also plays a role in how many atoms can fit

around a given molecule.

25

Note: This only applies to nonmetal atoms in Period 3 and later

Examples:

PCl5 Known to exist

NCl5 Not known to exist (N is too small for the to fit 5 Cl atoms around it)

Note: On homework problems only expand octets if there is no other way to

accommodate electrons or if the problem tells you to minimize the formal

charge.


Strength/Length of Covalent Bonds

Dissociation Energy (𝐷): The energy

required to separate bonded atoms

26

Avg. Bond Energies 𝒌𝑱

𝒎𝒐𝒍

H-H 432 O-H 467

H-F 565 O-O 146

H-Cl 427 F-F 154

C-H 413 Cl-Cl 239

C-C 347 O=O 495

C-N 305 C=O* 745

C-O 358 N≡N 941

N-H 391 C≡C 839

N-N 160 C≡N 891

*C=O (CO2) = 799



Double bonds

are not twice as

strong as a

single bonds.

Triple bonds are

not three times

as strong as a

single bond.

27


Student Question


Single bonds are ___________ triple bonds.

a) longer than

b) shorter than

c) the same length as

28

(ADDITION)


Student Question


Which of the following has the longest

carbon-oxygen bond?

Hint: You must draw the Lewis structures.

a) CO

b) CO32-

c) CO2

d) CH3OH

29


Take Away From Chapter 13

Big Idea: Bonds are formed from the attraction

between oppositely charged ions or by sharing

electrons. Only the valence electrons participate in

bonding. The shape of the molecules maximize the

distance between areas of high electron density.

Types of Bonding

Ionic Bonds (metal/non metal)

Be able to write electron configuration of ions. (26, 27, 29,

30)

Be able to predict size of ions.(23, 24, 25)

Be able to predict formula unit ionic compound.(33)

Covalent Bonds (non metal/non metal)

30

Numbers correspond to end of chapter questions.



Electronegativity Know the general electronegativity trend. (15)

Know that covalent bonds can have ionic because of

dipole moments Be able to identify the most polar bond. (16, 18, 19)

Know that ionic bonds can have covalent character

because of polarizablity

Lewis Structures

Be able to draw Lewis symbols (atoms).

Be able to draw Lewis structures of ionic compounds.

Be able to draw Lewis structures of covalent

compounds. (57, 58)

Know how to calculate formal charges.(78, 79)

Identification of most likely Lewis structure.

Know when multiple resonance structures are possible for a

compound. (60, 61, 65, 73)

Know when atoms can expand their octets (group 3 and

greater). (80)31





Know how to calculate Δ𝐻 from bond dissociation

energies .

△𝐻 = 𝐷 𝑏𝑟𝑜𝑘𝑒𝑛 − 𝐷 𝑓𝑜𝑟𝑚𝑒𝑑 (43, 44, 45, 46)

Know how to estimate the length of bonds.

Triple < Double < Single (74)

32


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